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JC 8819 



Bureau of Mines Information Circular/1980 



^70^/M^- 




Analytical Chemistry of the Citrate 
Process for Flue Gas Desulf urization 



By W. N. Marchant, S. L. May, W. W. Simpson, 
J. K. Winter, and H. R. Beard 



^^^NT^ 




UNITED STATES DEPARTMENT OF THE INTERIOR 



^.c 



Information Circular 8819 



Analytical Chemistry of the Citrate 
Process for Flue Gas Desulf urization 



By W. N. Marchant, S. L. May, W. W. Simpson, 
J. K. Winter, and H. R. Beard 




UNITED STATES DEPARTMENT OF THE INTERIOR 
Cecil D. Andrus, Secretary 

BUREAU OF MINES 

Lindsay D. Norman, Acting Director 








This publication has been cataloged as follows: 



United States. Bureau of Mines , 

Analytical chemistry of the citrate process for flue gas 
desulfurization. 

(Information circular - Bureau of Mines ; 8819) 

Bibliography: p„ 19. 

Supt. of Docs, no.: I 28.23:8819" 

L Flue gases— Desulphurization. 2. Citric acid. I. Marchant, 
Wayne N. II. Title. III. Series: United States. Bureau of Mines. Infor- 
mation circular ; 8819- 

TN295.U4 no. 8819 [TD885.5.S8] 622s 

[628.5'3'2] 79-607974 



CONTENTS 

Page 

Abstract 1 

Introduction 1 

Description of samples 3 

Analytical methods 3 

Sulfur species 3 

Hydrogen sulfide 3 

Polysulf ides 3 

Equipment and material 4 

Procedure 4 

Elemental sulfur 4 

Equipment and material 4 

Procedure 5 

Thiosulf ate ....;... 6 

Equipment and material . . ^ 6 

Procedure 6 

Poly thiona te s 7 

Equipment and material 7 

Procedure 8 

Dithionite 9 

Bisulfite 9 

Sulfate 10 

Equipment and material. 10 

^ Procedure 11 

Buffer species 12 

Citric acido 12 

Equipment and material 12 

Procedure 12 

Glycolic Acid .' 14 

Equipment and material 14 

Procedure 15 

Miscellaneous species 15 

Sodium 15 

Chloride 16 

General methods 16 

Direct-current polarography 16 

Equipment and material ^ 17 

Procedure 17 

Total sulfur 18 

Equipment and material 18 

Procedure 18 

Dis cus s ion 18 

References 19 

ILLUSTRATIONS 

1. Trichloroethylene solution absorbance as a function of sulfur 

concentration (1-millimeter cell) 5 



ii 



ILLUSTRATIONS —Continued 

Page 

2. Effect of trichloroethylene-to-aqueous volume ratio upon 

sulfur extraction 5 

3. Effect of contact time upon sulfur extraction 5 

4. Chromatogram of regenerated citrate absorbent; 0.25M 

thiosulf a te ; . 5M citric acid 7 

5. Chromatogram showing progress of citrate absorbent 

regeneration 9 

6 . Thin layer chromatographic detection of sulfur compounds 9 

7. Sample absorbance as a function of citric acid 

concentration 13 

8 . Percent reduction of apparent citric acid concentration 

caused by bisulfite and thiosulf ate 13 

9. Effect of pH upon apparent citric acid concentration 15 

10. Chromatogram of glycolic acid and thiosulf ate separation 

by HPLC 15 

TABLES 

1. Rf values for sulfur anions found in citrate process 

abs orbent 8 

2. Polarographic half -wave potentials and current per unit 

concentration for sulfur ions in citrate process 

solutions 17 



ANALYTICAL CHEMISTRY OF THE CITRATE PROCESS 
FOR FLUE GAS DESULFURIZATION 

by 

W. N. Marchant, ^ S. L. May, 2 W. W. Simpson, 1 J. K. Winter, 1 and H. R. Beard 1 



ABSTRACT 

The citrate process for flue gas desulfurization (FGD) is a product of 
continuing research by the U.S. Bureau of Mines to meet the goal of minimizing 
the objectionable effects of minerals industry operations upon the environment. 
The reduction of SO^ in solution by H^S to produce elemental sulfur by the 
citrate process is extremely complex and results in solutions that contain at 
least nine different sulfur species. Process solution analysis is essential 
to a clear understanding of process chemistry and its safe, efficient operation. 
The various chemical species , the approximate ranges of their concentrations 
in citrate process solutions, and the analytical methods evolved to determine 
them are hydrogen sulfide (~0M to 0.06M) by specific ion electrode, polysulfides 
(unknown) by ultraviolet (UV) spectrophotometry, elemental sulfur (~0M to 
r-^O.OOlM dissolved, ~0M to ^O.lll suspended) by UV spectrophotometry, thiosulfate 
(~0M to rJ).25M.) by iodometry or high performance liquid chromatography (HPLC), 
polythionates (~0M to r-O.OlM) by thin layer chromatogrphy (TLC), dithionite 
(searched for but not detected in process solutions) by polarography or TLC, 
bisulfite (^OM to 0.2M) by iodometry, sulfate (~0M to IM) by a Bureau-developed 
gravimetric procedure, citric acid (~(M to 0.5M) by titration or visible 
colorimetry, glycolic acid (~0M to IM) by HPLC, sodium (~1.5M) by flame photo- 
metry, and chloride by argent ome trie titration. 

INTRODUCTION 

One goal of the U.S. Bureau of Mines metallurgy research program is to 
reduce the objectionable impact of mineral-processing operations upon the environ- 
ment to the lowest possible level consistent with a viable minerals economy. 
The citrate process for flue gas desulfurization of (FGD) is a product of continuing 
Bureau research in pursuit of this goal (15, 17 ) ,^ and the analytical work 
described in this paper was developed to support citrate process research. The 

^Research chemist. Salt Lake City Research Center, Bureau of Mines, Salt Lake 

City, Utah. 
Research chemist (retired), formerly with the Salt Lake City Research Center, 

Bureau of Mines, Salt Lake City, Utah. 
^Underlined numbers in parentheses refer to the list of references at the end 

of this report. 



chemical reactions occurring in citrate process operation will be described 
in a subsequent paper. Briefly, the citrate process comprises the following 
four elements : 

1. SOg absorption. — Sulfur dioxide is absorbed from cooled and cleaned flue 

gas by contacting the gas with a solution of citric acid buffered at 
about pH 4.5. In this step, SO^ absorption lowers the solution pH to 
about 4. 

2. Absorbent regeneration .— High~purity elemental sulfur is precipitated by 

treating the SO^-rich absorbent solution with gaseous HgS. This restores 
the solution pH to 4.5, and permits the solution to be recycled for 
further absorption. 

3. Sulfur separation . — Elemental sulfur precipitated during the regeneration 

step is separated from the solution by flotation. It is freed from 
occluded or entrained absorbent solution by autoclave melting at about 
135" C. This permits withdrawal of pure molten sulfur from the lower 
phase in the autoclave. Supernatant absorbent from the upper phase, 
together with clarified absorbent from the flotation step, is then 
recycled for further SO^ absorption. 

4. HgS generation.— Approximately two-thirds of product sulfur is used for 

Jig b generation by reacting it catalytically at about 650" C with steam 
and a reducing agent such as natural gas, carbon monoxide, or methanol. 
The balance is cast into a form suitable for sale, storage, or shipment. 

Elemental sulfur formation from H^S and SOg in aqueous solution ("Wacken- 
roder's solution," after the German apothecary who first reported it) is a 
phenomenon known since 1845 (16) ; however, the chemistry of the process is 
extremely complicated and poorly understood. At least nine sulfur species are 
known to exist in solution during this reaction, and others have been postulated. 
Thus , citrate process research required the capability to analyze known sulfur 
compounds and to detect unknown compounds possibly present in process solutions. 
Methods for analyzing buffer species used in the process were also required, 
and are described herein. 

The complexity of sulfur chemistry encountered in this research was 
compounded by the inconsistent nomenclature used by different authors. For 
example, an aqueous solution of sulfur dioxide has been called variously 
sulfurous acid, sulfite, bisulfite, or simply SO^ . Similarly, the sulfur 
acid anions having the structure O3 S-S^^ -SOg^" have been called polythionic 
acids or polythionates most commonly; however, a prominent sulfur chemist recently 
advocated the name "polysulfanedisulfonate" (18) . By this convention, the 
compound O3 S-SSS-SOg^" , historically called pentathionate , would be called tri- 
sulfanedisulfonate. In the present paper each sulfur compound is identified 
by the name that most accurately describes its chemical form as it is encountered 
in the citrate process. Thus SO^ dissolved in a pH 4 to 4.5 buffered absorbent 
is called bisulfite (HSO3" ) rather than sulfite (SOg^"), which is essentially 
absent at this pH, or sulfurous acid, which is nonexistent. The traditional 
polythionate nomenclature is used rather than "polysulfanedisulfonate." 



While the latter name may be technically more descriptive, it is unwieldy 
and is generally unfamiliar. 

For easy reference, the paper is generally organized by compound. Sul- 
fur compounds are discussed in order of increasing formal oxidation state. 
Fore the sake of completeness, brief descriptions of published procedures 
used are included together with those developed by the Bureau during the 
research. 

DESCRIPTION OF SAMPLES 

Citrate process absorbent is typically 0.5M (molar) in citric acid. 
In addition, because thiosulfate rapidly accumulates in the absorbent during 
process operation, fresh absorbent is made 0.25M in sodium thiosulfate. Thus, 
other sample constituents must be determined against this high background con- 
centration of citric acid and thiosulfate. Recirculated absorbent is pale 
yellow and has a specific gravity of 1.1. Samples frequently contain ele- 
mental sulfur as a second phase, either as large floes or a finely divided 
suspension. These samples can usually be clarified by filtration, although 
it is often necessary to use membrane filters with nominal pore size less 
than 1 micrometer to accomplish this. 

Finally, samples typifying those from citrate process streams were known 
to contain at least the following solutes, at the concentrations shown where 
known: Sodium (1.5M), citric acid (0.5M) , thiosulfate (0.25M), polythionates 
(S^Og^" , with n = 3 to at least 6), sulfate (OM to ~1M) , hydrogen sulfide 
(OM to ^.06M), bisulfite (~0M to ^.2M), and elemental sulfur (~0M to ^O.lM). 

ANALYTICAL METHODS 

Sulfur Species 

Hydrogen Sulfide 

Hydrogen sulfide was analyzed with an Orion model 94-16* sulfide specific 
ion electrode and model 407 specific ion meter according to the manufacturer's 
instructions. Only mercury and silver are reported to interfere with this 
measurement, and no interference was observed from sample components during 
this research. Because samples often contain elemental sulfur, they were fil- 
tered through a glass fiber filter disk '(2 micrometers nominal pore size) prior 
to analysis. This precaution was taken to avoid accumulating sulfur on the 
electrode and not because the sulfur interfered with the analysis. The elec- 
trode is claimed by the maker to be sensitive to 10~'^M sulfide. It was used 
routinely for analysis of H^S from 10 parts per million to 2,000 parts per 
million in process solutions. ^ 

Polysulfides 

It is not believed possible to analyze individual polysulfide ions 
(HSj^~ or S„^~ ) in citrate process solutions because of the complex mixture 

^Reference to specific trade names or manufacturers does not imply endorsement 
by the Bureau of Mines . 



of other labile sulfur ions present. Several titrimetric methods have been 
published that use iodine or cyanide to qua.ntify polysulfides (19) , but these 
could not be expected to succeed because of the reactivity of thiosulfate , 
polythionates , and elemental sulfur toward these same reagents. To avoid the 
foregoing difficulty, spectrophotometry was used for sulfide analysis. 

Tetrasulfide and pentasulfide absorb light near 400 nanometers , with 
molar absorptivities , e, of 1,140 and 2,000, respectively (8^). The yellow 
color of recycled absorbent suggested that such ions might be present, but 
spectrophotometric scans of process solutions failed to reveal reproducible 
evidence for polysulfides. The method used would have detected tetrasulfide 
or pentasulfide at about 10" ^M. 

Equipment and Material 

A Gary model 15 UV -visible recording spectrophotometer was used to scan 
test solutions in fused quartz cells having either 1- or 10-millimeter path 
lengths . 

Procedure 

Samples were filtered to remove elemental sulfur and scanned from about 600 
nanometers to about 200 nanometers. In some samples, but not all, the spectrum 
showed an absorbance maximum at 388 nanometers, suggesting polysulfide absor- 
bance . Paradoxically, although all recycled citrate absorbent acquires a 
yellow color, only a few isolated test samples exhibited the absorbance at 388 
nanometers. Most samples showed no absorbance maximum in this region. The 
yellow color of these samples was due to a very weak (0.01 to 0.05 absorbance 
units for a 10-millimeter path cell) general absorbance of unknown origin 
extending into the visible region above 400 nanometers. 

Elemental Sulfur 

Moderate concentrations of solid elemental sulfur (greater than about 1 
to 5 grams per liter) were filtered and weighed; however, elemental sulfur is 
also present either dissolved or dispersed in unfilterable form. To analyze 
sulfur in these forms , a spectrophotometric method devised at the research 
center was used. This method uses as the basis for the determination the ultra- 
violet absorbance maximum exhibited at 268 nanometers by a trichloroethylene 
solution of sulfur. 

Equipment and Material 

Absorbance measurements were made on the Gary model 15 spectrophotometer. 
For the standard curve, 1-millimeter quartz cells were used. All other measure- 
ments were in 10-millimeter cells. Samples were shaken for extraction on a 
Burrell wrist-action shaker. Elemental sulfur (99.9 percent) was obtained from 
Montana Sulfur Go, Reagent-grade trichloroethylene (TGE) was obtained from 
Mallinckrodt. All other chemicals were the highest purity available commercially. 



Procedure 

The aqueous sample solution (10.0 milliliters) was contacted in a 60-milli- 
liter separatory funnel with 10.0 milliliters of TCE by shaking for 20 minutes 
on the wrist-action shaker. After phase disengagement, the optical absorbance 
of the TCE solution was measured at 268 nanometers. Figure 1 shows the absor- 
bance of solutions containing varying amounts of pure elemental sulfur. The 
molar absorptivity of sulfur under these conditions is 850. 




ELEMENTAL SULFUR CONCENTRATION, millimolar 

FIGURE 1> - Trichloroethytene solution absorbance as a func- 
tion of sulfur concentration (1-millimeter cell). 



In order to establish 
optimttm conditions for the 
analysis , sulfur extraction 
was determined as a function 
of phase ratio and contact 
time as shown in figures 2 
and 3, respectively. As may 
be seen from these figures , 
the small incremental improve- 
ment in sulfur extraction 
obtainable by increasing the 
TCE -to-aqueous ratio above 1 
would be offset by an atten- 
dant decrease in sensitivity 
for the more dilute TCE solu- 
tion. Similarly, no signif- 
icant benefit would be gained 
by a contact time greater than 
20 minutes . 




0.5 1.0 1.5 2.0 

TCE-to-AQUEOUS VOLUME PHASE RATIO 
FIGURE 2» • Effect of trichloroethylene-to-aqueous 
volume ratio upon sulfur extraction. 




FIGURE 3. 



5 10 15 20 

CONTACT TIME, minutes 

Effect of contact time upon sulfur 
extraction (1:1 phase ratio). 



25 



Tests were run to determine whether other ions present in typical citrate 
process samples would interfere with this analysis. No interference was 
observed for solutions 0.5M in citric acid and also containing 0,2M sulfide, 
thiosulfate, bisulfite, trithionate , tetrathionate , or dithionite. 

The sulfur concentration in the original aqueous sample is calculated 
from the equation, 

A , 
C = 850 b 

where C = sulfur concentration in moles per liter , 

A = solution absorbance in absorbance units , 

and b = cell path length in centimeters. 

Samples for which the absorbance exceeds the instriraient scale may be 
diluted with TCE without affecting the determination, in which case a correction 
for dilution must be applied to the relationship above. 

Thiosulfate 

Two methods were used for thiosulfate determination. In most cases stan- 
dard iodometric titration was suitable (i6 ) . (See also "Bisulfite" section of 
the present report). Interference by bisulfite was prevented by masking with 
formaldehyde; a correction to the iodine titer must be applied for hydrogen 
sulfide present in the sample. 

During investigations of citrate process chemistry and reactions , a thio- 
sulfate analysis was required that was not susceptible to error from unknown 
components of sample solutions that might consume iodine, thereby introducing 
an indeterminate error in the iodine titer. A high-performance liquid chromato- 
graphic (HPLC) procedure was developed to supplement the iodometric method. 

Equipment and Material 

A Perkin-Elmer model 1220 liquid chrcmatograph and model LC55 variable 
wavelength UV detector coupled to a Beckman model 1005 strip chart recorder were 
used for the analysis. The sample Injection valve was fitted with a 30- 
microliter sample loop. The analytical colirain was 2 .6-millimeter-inside-diameter 
by .5-meter-long stainless steel packed with Zipax strong anion exchange 
pellicular resin (DuPont Instrument Co.). Peak areas were integrated elec- 
tronically by a Spectra- Physics Autolab System IV data analyzer. Reagent-grade 
cadmium sulfate in degassed distilled water was used as the eluent (mobile 
phase) . 

Procedure 

Prior to injection onto the column, samples were diluted fiftyfold with 
distilled water to prevent overloading the low-capacity resin. Samples were 



eluted with a solution of 0.38 gram cadmiinn sulfate per liter of degassed 
distilled water at a flow rate of 1.5 milliliters per minute (1,000 pounds 
per square inch pressure). Figure 4 shows the analysis of regenerated absor- 
bent solution containing 0.25M thiosulfate in 0.5M citric acid buffer. The 
significant feature of this analysis is that thiosulfate, which elutes first 
among the components of typical process solutions, is separated from citric 
acid. By all other HPLC procedures screened during development of this pro- 
cedure, citric acid and thiosulfate eluted together, making thiosulfate 
determination impossible. The thiosulfate retention time of 7.7 minutes is 
reproducible to within ±0.3 minute over the range from O.lM to 0.25M, 

Polythionates 



2- 



Polythionates , S„ Og with n^3 , are always present in citrate process 



<n 

z 
o 

Q. 
CO 
UJ 

q: 

UJ 

o 
oc 
o 
o 

UJ 



Thiosulfate 



solutions. Dithionate , S^Og^", which is structurally similar, is chemically 
distinct from higher members of the series and was not detected in process 
solutions. The detection limit for dithionate by methods described herein is 
estimated to be 50 parts per million. Although several methods have been 
reported for analyzing individual polythionates, none is satisfactory for 
large numbers of samples . Kurtenacker and Goldbach ( 12 ) determined polythionates 
through Sg0g^~ in mixtures by titrating five aliquots of each sample. Jay 
(9.) published a procedure for determining total but not individual ^polythionates 
in a sample; however, this method was not reliable in the present study, 

possibly because of interference from 
the much higher concentrations of 
other solutes. Chapman and Beard (1^) 
reported a HPLC method for individual 
polythionates , but this method requires 
high sample dilution, and polythionates 
are not detected below about 1 to 2 
grams per liter in the original sample. 
Polythionate analysis by thin-layer 
chromatography (TLC) was found to be 
a rapid and convenient procedure com- 
patible with citrate process research 
needs , 

Equipment and Material 

High-surface-area silica gel plates 
(0.2 millimeter thick) equivalent to 
Merck 5775 are most satisfactory. A 
Kotes System I densitometer with baseline 
corrector was used for quantifying 
developed plates. A Soltec model 281 
strip-chart recorder was used to con- 
vert densitometer output to optical 
density spectra. Peaks were integrated 
with the Spectra-Physics Systems IV 
data analyzer. Reagent-grade chemicals 




30 



10 20 

TIME, minutes 

FIGURE 4« - Chromatogram of regenerated citrate 
absorbent; 0.25M thiosulfate; 0.5M 
citric acid. 



were used in distilled, deionized water to prepare eluent solutions. Potassium 
trithionate and pentathionate were synthesized by the method of Goehring, Statnm, 
and Feldman (4-5). Sodium tetrathionate was purchased from ICN Pharmaceuticals, 
Inc. 

Procedure 

Samples were diluted 1:50 with distilled water prior to chromatography. 
Those samples containing solid sulfur were filtered prior to analysis to avoid 
plugging the spotting syringe. Chromatograms were developed by ascending 
solvent consisting of ethanol, £-dioxane , concentrated ammonia, water (12:4:2:1) 
Developed chromatograms were visualized by spraying with either ethanolic 
silver nitrate to aqueous palladium chloride acidified with a few drops of HCl 
in 50 milliliters of solution. Silver nitrate is the more sensitive reagent; 
however, plates sprayed with silver nitrate darken after 1 or 2 hours and are 
unsuitable for densitometry. Plates visualized with palladium chloride remain 
essentially unchanged and may be quantified at any time after spraying. Figure 
5 illustrates the TLC analysis of SO^ -loaded absorbent during regeneration 
with H^S. This plate was visualized with silver nitrate. Figure 6 shows the 
separation and corresponding densitometer scan of 3 microliters of a synthetic 
mixture containing thiosulfate (spot 1, 7.99 grams per liter), trithionate 
(spot 2, 3.58 grams per liter), tetrathionate (spot 3, 1.26 grams per liter), 
and pentathionate (spot 4, 0.77 gram per liter). The silica gel plate was a 
commercial product on a transparent plastic backing that permitted operation of 
the densitometer in the visible transmission mode. Spot 4 corresponding to 
pentathionate was too faint to photograph satisfactorily; nevertheless, quan- 
tification was possible. By this method, polythionate concentrations at least 
as low as 100 parts per million may be analyzed. Table 1 lists the R^ values 
for thiosulfate and the polythionates analyzed in process solution by this 
method .5 

TABLE 1, - Rf values for sulfur anions found 
in citrate process absorbent 



Compound^ 


R, 


s^o,^- 


0.06 


2 "3 

S, ^~ 


.44 


3 '-'S „ 

s ^- 


.61 


4 G 3 

Sk Os .". 


.72 







^Dithionate, S^Og^" , elutes between thio- 
sulfate and trithionate in this solvent 
system but was never detected in citrate 
process solution. 



^Rf is a dimensionless parameter that indicates the relative distance that a 
given spot has traveled from the origin. It is defined as the ratio of 
distances from the origin to the center of the spot and from the origin to 
the solvent front. Thus, a spot which does not move from the origin has an 
R^ = 0; a spot which moves with the solvent front has a R^ = 1.00. 



* • ♦ 



* * • ^ 



ss '4 * 



* *# *^ ##(►###•* 



If llllllllflllllllil I 



Elemental sulfur 
Higher polythionates 



S4O6' 



^ 



S3O6 



2- 



Thjosulfate 



50 100 150 200 

ELAPSED TIME IN CONTACT WITH HgS, minutes 

FIGURE 5t - Chromatogram showing progress of citrate absorbent regenerationi 



A 



Dithionite 



2-1 



2 (SaOg' ) 



u 



3 (8406'^") 



ooiveni 
front -»• 


4 






3 


* 




2 


• 


Origin* 


1 


'- 




TLC plate 



4 (SgOg^-) 



FIGURE 6t - Thin layer chromatographic detec- 
tion of sulfur compounds. 



;3 + 



Dithionite (S^O^^"), a compound of 
was regarded as a possible inter- 
mediate in the reduction of SO^ in the 
citrate process. It is not separated 
from the known components of process 
solutions by the TLC solvent system 
above; however, a solvent consisting 
of methanol, n-butanol, water (3:1:1) 
separates dithionite and thiosulfate 
with Rf values of 0.80 and 0.64, 
respectively. Dithionite may also be 
visualized by spraying with either 
ethanolic silver nitrate or aqueous 
palladium chloride. Extensive exami- 
nation of process solutions failed to 
find dithionite under conditions that 
should have detected about 50 parts 
per million or less. 

Bisulfite 

Bisulfite (HSOg") was determined 
by a standard iodometric method. Because 
thiosulfate interferes with this titra- 
tion, it is necessary to titrate two 
aliquots of sample , one of which is 
pretreated with formaldehyde to seque- 
ster bisulfite as its addition compound 
with formaldehyde (6^, pp. 223-231), 



10 



A suitable aliquot of sample (2 to 10 milliliters) is added to about 50 
milliliters of water. Formaldehyde (5 milliliters of 37-percent-HCHO solution) 
is added, followed by sufficient NaOH (~1M) to make the solution aklaline to 
phenolphthalein. The solution is allowed to stand for 10 minutes to allow 
complete reaction of dissolved SOg with the formaldehyde. 

A second aliquot of sample (2 to 10 milliliters) is added in one portion 
to an excess of standard iodine solution (-jO.IN) contained in about 50 milli- 
liters of water. The volume of iodine solution is recorded as "A." The sample 
is acidified with 2 milliliters of glacial acetic acid and titrated with 
standard thiosulfate solution (^.IN) to a faint yellow color. Starch indi- 
cator is then added, and titration with the standard thiosulfate solution is 
continued to a colorless end point. The total volume of thiosulfate is recorded 
as "B." 

Titration of the first sample aliquot is now completed by adding 2 milli- 
liters of glacial acetic acid to the solution, followed by starch indicator and 
then titrating with standard iodine solution (^,1N) to the blue end point. The 
volume of standard iodine solution is recorded as "C." 

Analyte concentrations, in grams per liter, are calculated as follows: 

[(A)(Nl3)-(B)(NS3 03)-(C)(Nl3)]32. 



bisulfite concentration = 



thiosulfate concentration = 



milliliter of sample 

(C)(NI^)(112) 



milliliter of sample 

where Nig = normality of standard iodine solution, 

and NS2O3 = normality of standard thiosulfate solution. 

Iodine solution is standardized against arsenious oxide (As^Og) as either 
the dry reagent or standard solutions available from chemical suppliers. The 
thiosulfate solution is then standardized against the standard iodine. 

Sulfate 

A gravimetric method developed by the Kraft paper industry was adapted by 
the Bureau for use with citrate process samples. Although simplified, it does 
not sacrifice accuracy for samples containing from 0.05 to 50 grams per liter 
of sulfate, which represents the range of concentrations encountered. 

Equipment and Material 

An lEC model K centrifuge was used for separating the first BaSO^ preci- 
pitate. Carbon dioxide used for deoxygenating solutions was obtained in cylin- 
ders from commercial vendors. All chemicals were reagent grade. Iodine 



11 



reagent (about 0.3N) for selectively oxidizing thiosulfate to tetrathionate 
was prepared by dissolving 38.1 grains iodine and 120 grams potassium oidide 
in about 50 milliliters water and diluting to 1 liter after complete disso- 
lution. Distilled water (33 to 500 milliliters) was deoxygenated by bubbling 
COg through it for 5 minutes and was then stored in a tightly stoppered bottle. 
Two-percent solutions of HCl and NaOH were deoxygenated in the same way. 
(Some carbonate forms in the NaOH solution; however, this does not interfere 
with the analysis.) 

Procedure 

The analysis is based on the classical precipitation of sulfate as the 
barium salt; however, because of the high concentrations of solutes other than 
sulfate in process samples and the low solubility of other barium salts such 
as citrate, samples were analyzed by the following four-step method: 

1. Choose an aliquot of sample to contain 1 to 125 milligrams of sulfate. 
Pipet the aliquot into a 150-milliliter beaker and add 5.0 milliliters of 37- 
percent-formaldehyde solution and 5.0 milliliters of isopropyl alcohol. 
Dilute to 40 milliliters with deoxygenated water and adjust the pH to 4.0 
with deoxygenated HCl or NaOH as required. Allow the sample to stand for 5 
minutes to insure complete reaction of bisulfite and formaldehyde. (This 
will prevent oxidation of bisulfite to sulfate during addition of iodine in 
the next step.) During this interval, purge oxygen from the 0.3N iodine 
solution by bubbling CO^ through it. Add iodine reagent to the sample to a 
blue starch indicator end point, then add 1 milliliter excess. This converts 
thiosulfate quantitatively to tetrathionate. 

2. Add about 5 grams of lump-free ammonium chloride to a 250-milliliter 
centrifuge bottle, displace air in the bottle with CO^ , and add the sample from 
step 1, washing the sample from the beaker with deoxygenated water. Adjust 
sample volume to about 175 milliliters. Swirl the stoppered centrifuge bottle 
until the ammonium chloride dissolves, then add 1.25 to 1.5 grams solid 

BaClg '21^ and swirl thoroughly. Centrifuge at 2,000 revolutions per minute 
for 30 minutes. Repeat if necessary to clarify the supernatant. When a clear 
supernatant is obtained, decant and discard it. 

3. Transfer the precipitate from step 2 quantitatively into a 400-milli- 
liter beaker. Dilute with water to a volume of about 400 milliliters, and add 
2 milliliters of 2-percent HCl. Heat on a hotplate at medium heat for 30 
minutes, then add 10 milliliters concentrated HCl, cover with a watch glass 
and boil for 2 minutes . Remove the beaker from the heat , add methyl orange 
indicator, neutralize with cencentrated aqueous ammonia, reacidify with 10 
drops of excess concentrated HCl, and boil for 5 minutes. Remove from heat 
and cool overnight. Filter the supernatant solution through filter paper 
equivalent to S. and S. No. 589 containing 5 milliliters pulp. (Do NOT trans- 
fer the bulk precipitate to the filter paper.) Wash down the sides of the 
beaker containing the BaSO precipitate with 20 milliliters concentrated nitric 
acid and 20 milliliters concentrated perchloric acid. Place the filter paper 
and pulp into this beaker and digest to heavy fumes over medivmi heat. 



12 



If the residue is colored, add concentrated nitric acid and continue fuming 
until colorless. Finally, fvm.e down to a 2-milliliter volume, cool, and 
add 250 milliliters distilled water, methyl orange indicator, neutralize 
with aqueous ammonia, and reacidify with 6 drops excess concentrated HCl. 

4. Heat the sample, add 5 milliliters of a solution of 225 grams 
BaClg* 2Hg in 2 liters of water. Cover the sample and boil for 5 minutes, 
then allow to simmer for 30 minutes. Remove from heat and cool overnight. The 
precipitate can now be recovered by filtration, after which BaSO^ is determined 
by heating the filter paper plus precipitate at 900° C for 30 minutes in an 
annealing cup and weighing the BaSO^ . 

Buffer Species 

Citric Acid 

For process control and material balance calculations, it is essential to 
be able to quantify the buffer component of process solutions as well as the 
reactants . The following procedures were used. 

Equipment and Material 

Automatic titrations were performed on a Mettler DKlO-DKll titrimeter with 
model GAIO recorder. Visible absorbance was measured on the Cary model 15 
spectrophotometer. All chemicals were reagent grade. 

Procedure 

A Bureau-developed titrimetric method was used for most laboratory experi- 
ments and for pilot plant testing under this program. The basis for the method 
is the release by ordinarily tribasic citric acid of a fourth acidic hydrogen 
upon coordination with cupric ion (20). A 1.0-milliliter aliquot of sample is 
diluted to about 50 milliliters with water. Hydrogen peroxide (5 milliliters 
of 30-percent solution) is added, and the solution is heated to 90° C for 5 
minutes to destroy bisulfite, which interferes with the titration. After 
cooling, the sample is titrated to just past the end point with standarized 
NaOH solution, about O.lN. (The titrimeter should be operated in the differen- 
tial mode to provide sharp definition of the end points.) Copper sulfate penta- 
hydrate , about 0.6 gram, is then added, and the sample is titrated to a second 
end point. The volume of titrant added between the two endpoints is recorded. 
At the end of the titration the concentrations of citrate and cupric ion should 
be approximately 0.005M-and 0.02M, respectively. The citric acid concentration 
(C) , in grams per liter is calculated from 

C = (N)(V) (192.1), 

where N = normality of NaOH titrant, 

and V = volume of titrant added between end points, in milliliters. 



13 



An alternate citric acid analysis used by the dairy industry (T) has 
recently been found to be suitable for citrate process samples and to require 
less instrumentation and operator skill. In this procedure, 1.0 milliliter of 
sample containing not more than 1 micromole (192 micrograms) of citric acid is 
placed in a test tube, followed by 1.3 milliliters of pyridine. The solution 
is mixed by swirling. Acetic anhydride (5.7 milliliters) is added, after which 
the contents are swirled briefly, then incubated for 30 minutes in a constant 
temperature bath at 32° ±0.25** C. Citrate acid is then determined by measuring 
the absorbance of the resulting solution at 435 nanometers and comparing sample 
absorbance with a standard curve. 

Because the author of this method stated that solution absorbance in this 
analysis does not obey Beer's law exactly, absorbance was measured for a series 
of samples between 0.2 millimolar and 2 millimolar. Figure 7 shows absorbance 
as a function of citric acid concentration in this range . While the curve is 
not liniar, it is nearly so up to 1 millimolar, and satisfactory values are 
possible by comparing sample absorbance with such a standard curve even beyond 
1 millimolar. 



To determine whether other potential components of sample solutions would 
interfere with this analysis, it was repeated for solutions containing 0.02M 
Fe2+, Cu^+, Ni^+, Cr3+, Sn^^ , Ca^+ , Mg^+ , Al^^ , Zn^+ , Na+ , CI", or SO^^" in 
O.IM 




0.4 0.8 1.2 1.6 2.0 
CITRIC ACID CONCENTRAHON, millimolar 

FIGURE 7. - Sample absorbance as a function 
of citric acid concentration. 




0.1 0.2 0.3 0.4 
CONCENTRATION, molar 

FIGURE 8. - Percent reduction of apparent citric 
acid concentration caused by bi- 
sulfite and thiosulfate. 



0.5 



14 



citric acid. No interference was seen; however, the following precautions must 
be observed when using this analysis: 

• Both bisulfite and thiosulfate depress sample absorbance. Figure 8 
shows the percent by which the apparent concentration of known 0.5M 
citric acid is reduced by varying amounts of these solutes. Because 
this effect is linear as shown, a correction may be applied to the 
apparent citric acid analysis when these components are present in 

the sample. A correction curve such as figure 8 should be prepared for 
each system under study; nevertheless, error due to bisulfite or thio- 
sulfate is not believed to be a serious impediment to use of this 
procedure. At the greatest concentrations of these solutes usually 
encountered, 0.15M and 0. 25M for bisulfite and thiosulfate, respectively, 
the corresponding errors are 4 percent and 2.3 percent. Furthermore, 
only absorbent freshly contacted with SO^ contains such a high con- 
centration of bisulfite. At all other points in the system the bisul- 
fite concentration is below about 2 grams per liter, at which level 
the corresponding error is below 1 percent. 

• Samples must be diluted carefully in volvimetric glassware to reduce 
the citric acid concentration to about O.OOlil prior to adding pyridine. 

• Accurate citric acid results are obtained for samples with pH below 
4.0; however, above pH 4.0, the apparent concentration is low as 
shown in figure 9. This effect is easily prevented by adjusting 
sample pH to 4 or lower prior to diluting and adding pyridine. 

• Both pyridine and acetic anhydride are degraded by light and/or 
moisture. Fresh reagent should be used daily, and stock reagent should 
be stored dry and in low actinic glass containers. 

Glycolic Acid 

Glycolic acid (hydroxyacetic acid) was successfully tested as a possible 
lower cost alternative to citric acid buffer in the citrate process (15) . It 
may be analyzed as the methyl ester by gas chromatography; however, this is a 
relatively time-consuming procedure. A simple HPLC method has been found to be 
suitable. Although no exhaustive work has been done on this analysis, the 
preliminary results reported here show that HPLC can be used to separate and 
quantify glycolic acid in citrate process absorbent. 

Equipment and Material 

The liquid chromatograph and associated equipment were as described for 
thiosulfate analysis. The analytical column was 2.6-millimeter-inside-diameter 
by 0.5 -meter- long stainless steel cleaned and packed as described by Kirkland 
(10) with Permaphase AAX weak anion exchange resin (DuPont). 



15 





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3 4 5 6 

PH 
Effect of pH upon apparent citric 
acid concentration. 



LU 

(/) 

z 
o 

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V) 
UJ 
IT 

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o 
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50 



FIGURE 10. 



100 150 

TIME, seconds 



200 



250 



Chromatogram of glycolic acid and 
thiosulfate separation by HPLC. 



continuous internal lithium standard. 
1 part per million. 



Procedure 

A 30-microliter sample loop was 
used to apply samples to the column. 
Elution was by reagent-grade sodium 
sulfate (7 x 10"^ M) in degassed 
distilled water at a flow rate of 1 
milliliter per minute (1,000 pounds 
per square inch). Sample components 
were detected by UV absorbance at 205 
nanometers . 

Figure 10 shows the analysis 
obtained in a l.OM solution of glycolic 
acid containing 0.6M sodium thiosulfate. 
As with citric acid, many eluents 
failed to separate the glycolic acid 
buffer from thiosulfate, which is always 
found in process solutions. Sodium 
sulfate eluent provided good resolution 
of the glycolic acid and thiosulfate 
peaks . The method is limited by the low 
UV absorptivity of glycolic acid; how- 
ever, it has been used for solutions as 
low as 0.05M. To avoid overloading the 
low-capicity resin, samples were diluted 
with water prior to injection. For 
the analysis of figure 10, a 500-fold 
dilution was used. Less dilution may 
be used to improve the sensitivity for 
glycolic acid if the concentrations of 
other solutes are found to be lower than 
the 0.6M thiosulfate in the analyte of 
figure 10 . 

Miscellaneous Species 

Sodium 

Sodium ion was analyzed by flame 
photometry. An Instriiment Laboratory 
Inc. model 343 photometer was used with 
The limit of detection by this method is 



Undiluted citrate process solutions are sufficiently viscous that pre- 
cautions must be taken to insure that the sample aspiration rate is comparable 
to that of the sodium standard solutions. This may be accomplished by either 
(1) diluting the sample prior to analysis, if the sodium concentration after 
dilution remains above the detection limit (process samples were routinely 



16 



diluted 100:1 with water) or (2) adding citrate to standard solutions at a 
concentration equal to that of the sairple. 

Chloride 

Process solutions were analyzed for chloride ion by an adaptation of the 
Volhard method. Because sulfide and sulfite interfere with the analysis, a 
preliminary treatment with H^O^ is necessary to oxidize sulfite to sulfate 
followed by boiling to decompose excess H^Qg and to expel residual sulfide. 

In a typical analysis, an aliquot of sample containing about 5 to 100 
milligrams of chloride is placed in a Erlenmeyer flask and diluted with water 
to about 100 milliliters. About 1 milliliter of 30-percent -H^ Og solution is 
added, after which the flask is covered with a watch glass and boiled for about 
5 minutes. After cooling, an excess of standard AgNOg (about O.lN) is added; 
then, the sample is acidified with 2 milliliters of concentrated nitric acid. 
Iron(III) indicator (2 milliliters of saturated FeNH^ (SO^ )^ solution) is added 
followed by 5 milliliters of chloride-free nitrobenzene. After the flask is 
vigoriously shaken, the excess AgNOg is titrated with standard (O.lN) potassium 
thiocyanate solution to the first red FeSCN^''' color that persists for 1 minute. 

The chloride concentration in grams per liter, is then calculated from 

Cr - r (A) (B)-(C) (0)1(35.46) , 

E 

where A = milliliter of standard AgNOg added, 

B = normality of standard AgNOg , 

C = milliliter of standard potassium thiocyanate used, 

D = normality of standard potassium thiocyanate, 

and E = milliliter of sample taken. 

The foregoing procedure is satisfactory for solutions containing more than about 
100 parts per million of chloride. In solutions containing less than 100 parts 
per million of chloride, a chloride-specific ion electrode is recommended. 

General Methods 

Direct-Current Polarography 

One of the principal objectives of citrate process laboratory research 
was to identify the compounds of sulfur that exist in an operating system. In 
addition to TCL, direct-current (DC) polarography was used to screen process 
samples for previously unknown compounds. 



17 



Other investigators have used polarography to analyze thiosulfate , sulfide, 
trithionate , dithionite , and bisulfite (3_, 11-13), but the supporting elec- 
trolyte in these investigations was phosphate, acetate, nitrate, or chloride. 
Because samples generated in citrate process research are in relatively con- 
centrated (0.5M) citric acid, this was used as supporting electrolyte for these 
tests, diluted to a concentration in the cell of O.OIM. 

Equipment and Material 

The polarographic cell consisted of a 1-liter, five-neck, round flask. 
The dropping mercury electrode was inserted in the central neck, nitrogen for 
deaerating the solution was admitted through the rear neck, the front neck accomo- 
dated the saturated calomel reference electrode, and the necks on either side 
of center were for two graphite auxiliary ("working") electrodes. Potential to 
the cell was supplied by a Princeton Applied Research model 173 potentiostat 
and was varied by an Elscint model ABA-26 automatic baseline advance. Applied 
potential was displayed on a Hewlett-Packard model 3440A digital voltmeter. 
Current, in microamperes, was displayed on the potentiostat and recorded simul- 
taneously on a Honejrwell Electronic 19 recorder. The dropping mercury elec- 
trode was made from standard capillary bore glass tubing and was regulated to 
deliver 1 drop every 3 seconds. 

Procedure 

Ten milliliters of sample were diluted with distilled water to 1.0 liter 
in the cell, after which the solution was deaerated by bubbling nitrogen 
through it for 10 minutes . Because the half-wave potentials of dithionite and 
sulfide were indistinguishable below pH 7, the sample pH was adjusted to 7 
prior to final dilution. At this pH, the known sulfur ions could be analyzed 
in mixtures such as process solutions. The solution was then analyzed by 
scanning between +350 millivolts and -1,000 millivolts at a rate of 200 milli- 
volts per minute. Table 2 lists the half-wave potentials and corresponding 
current measured for electroactive sulfur ions during these tests. 

TABLE 2. - Polarographic half-wave potentials and current 
per unit concentration for sulfur ions in 
citrate process solutions 



Ion 


Half-wave^ potential , 
millivolts 


Current per millimole , 
microamperes 


s-^- 


-620 
-150 
-500 
+250 


1.5 


hK^ 


4.5 
1.25 


2 4 

HSO„ 


5.5 


3 





^Referred to the saturated calomel electrode, 



No polarographic wave was observed for sulfate or the polythionates . 



18 



The main objective of the polarographic analyses was not to quantify the 
sulfur ions, which was done by other methods described, but to search for com- 
pounds not previously known; however, no reproducible evidence for the existence 
of previously undetected compounds was obtained. 

Total Sulfur 

The total sulfur content of test solutions was required both as a check of 
other analyses and for material balance calculations. It was determined by 
oxidizing all sulfur to sulfate, which was then determined gravimetrically as 
the barivtm salt. 

Equipment and Material 

Only standard laboratory glassware was required in addition to a magnetic 
stirrer-hotplate. For several simultaneous analyses a 230-volt hotplate having 
a 45- by 60-centimeter surface was used. 

Procedure 

Aliquots of sample (5.00 milliliters) in 250-milliliter beakers were diluted 
to about 50 milliliters with water. Bromine (about 2 milliliters) was added after 
which the sample was heated with magnetic stirring until no bromine droplets 
were visible. Concentrated nitric acid (5 milliliters) was added, and the solu- 
tion was then boiled gently to a straw yellow color. Concentrated HCl (20 milli- 
liters) was then added, and boiling was maintained to expel nitrates. Further 
20-milliliter additions of HCl were made with continued boiling until the solu- 
tion was colorless. The solution was then concentrated by boiling to a volume 
of about 25 to 50 milliliters, after which 15 milliliters of 0.25M BaClg solu- 
tion was added. The solution was boiled gently for 10 minutes, allowed to 
cool, and filtered under gentle suction onto previously tared Whatman GF/C glass 
fiber disks . For routine purposes , predrying of the filter disks was not justified 
as they were found to contain less than 0.15 percent moisture by weight. 

This relatively rapid and simple analysis produced replicate sulfur values 
typically within 2 percent or less. 

DISCUSSION 

The methods described herein comprise those analytical procedures that 
were evolved during the 7-year period in which the citrate process developed from 
preliminary laboratory experiments to a demonstration plant under construction 
to treat exit gas from a 60-megawatt coal-buring powerplant. They should pro- 
vide satisfactory resul-ts for other investigators studying sulfur chemistry as 
encountered in FGD systems, particularly of the regenerable type. 

During preparation of this manuscript, an innovation infrared spectrophoto- 
metric method for determining citric acid, thiosulfate , and sulfate simultaneously 
in scrubber solutions was published (14) . This method was not studied during the 
present investigations, but it may be of interest to other researchers having 
access to a modern double-beam infrared spectrophotometer. 



19 



REFERENCES 

1. Chapman, J, N. , and H. R. Beard. Fast Analytical Procedure for 

Separation and Determination of the Polythionates Found in Wacken- 
roder's Solution by High Speed Liquid Chromatography. Anal. Chem. , 
V. 45, 1973, p. 2268. 

2. Furness , W. J. The Analysis of Commercial Dithionates and Related 

Compounds by Polarographic Methods. Soc, Colourists and Dyers, 
V. 66, 1950, pp. 270-277. 

3. Garber, R. W. , and E. W. Wilson. Determination of Atmospheric Sul- 

fur Dioxide by Differential Pulse Polarography . Anal. Chem., v. 
44, 1972, pp. 1357-1360. 

4. Goehring, M. , and U. Feldmann. Preparation of Potassium Penta- and 

Hexathionates , Z. Anorg. Chem., v. 257, 1948, pp. 223-226. 

5. Goehring, M. , H. Stamm, and U. Feldmann. Polythionic Acids and Their 

Formation. Z. Anorg. Allg. Chem., v. 250, 1942, pp. 56-71. 

6. Haff, L. V. The Analytical Chemistry of Sulfur and Its Compounds. 

Wiley Interscience , New York, 1970, p. 235. 

7. Hartford, C. G. Rapid Spec trophotome trie Method for the Determination 

of Itaconic, Citric, Aconitic , and Fumeric Acids. Anal. Chem., 
V. 34, 1962, p. 426. 

8. Hoffman, M. R. Kinetics and Mechanism of Oxidation of Hydrogen Sulfide 

by Hydrogen Peroxide in Acidic Solution. Environ. Sci. and Tech., 
V. 66, 1977, p. 61. 

9. Jay, R. R. Determination of Polythionic Acids. Anal. Chem., v. 25, 

1953, p. 288. 

10. Kirkland, J. J. Modern Practice of Liquid Chromatography. Wiley 

Interscience, New York, 1971, p. 178. 

11. Kolthoff, I. M. , and C. S. Miller. The Reduction of Sulfurous Acid 

(Sulfur Dioxide) at the Dropping-Mercury Electrode. J. Amer. Chem. 
Soc, V. 63, 1941, pp. 2818-2821. 

12. Kurtenacker, A., and E. Goldbach. The Analysis of Polythionate Solu- 

tions. Z. Anorg. All. Chem., v. 166, 1927, pp. 177-189. 

13. Lem, W. J., and M. Wayman. Decomposition of Aqueous Dithionite. Can. 

J. Chem., V. 48, 1970, pp. 776-781. 



■^wi^il^^l^»^^l^l^ll^ ^ ^lll | l!!llll ) l ^| !ll ! l|l | |l ^ l lwl l ^| l! ll^l!,l |■ 



20 



14. Luke, K. W., and L. 0. Edwards. Application of Infrared Spectro- 

photometry to the Analysis of Citrate, Thiosulf ate , and Sulfate 
in Citrate Scrubber Solutions. Radian Corp. Tech. Note 200-203-02, 
June 1978, 34 pp., available for consultation at U.S. Bureau of 
Mines, Salt Lake City Research Center, Salt Lake City, Utah. 

15. Nissen, W. I., L. Crocker, and D. A. Martin. Lead Smelter Flue 

Gas Desulfurization by the Citrate Process. Ch. 52 in World 
Mining and Metals Technology, ed. by Alfred Weiss. (Proc. Joint 
Min. and Met. Inst, of Japan-AIME Meeting, Denver, Colo., Sept. 
1-3, 1976), Port City Press, Baltimore, Md., v. 2, 1976, pp. 
825-854. 

16. Roscoe , H. E., and C. Schorlemmer. A Treatise on Chemistry. 

MacMillan and Co., Ltd., London, 1920, p. 463. 

17. Rosenbaum, J. B., W. A. McKinney, H. R. Beard, L. Crocker, and 

W. I. Nissen. Sulfur Dioxide Emission Control by Hydrogen Sulfide 
Reaction in Aqueous Solution. The Citrate System. BiiMines 
RI 7774, 1973, 31 pp. 

18. Schmidt, M. Sulfur in Organic and Inorganic Chemistry. Marcel Dekker, 

Inc., New York, 1972, p. 96. 

19. Szekeres , L. Analytical Chemistry of the Sulfur Acids. Talanta, 

v. 21, 1974, pp. 1-44. 

20. Warner, R. C, and I. Weber. The Cupric and Ferric Citrate Complexes. 

J. Amer. Chem. Soc, v. 75, 1953, p. 5086. 



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